EL2

Spectroscopy

Under certain conditions, a substance will absorb or emit electromagnetic radiation in a characteristic way. Analysis of this can lead to identification of the substance.

Absorption Spectra

Glowing stars emit all frequencies between UV and IR

Chromosphere contains ions, atoms, and small molecules

These absorb radiations

Emitted light is missing specific frequencies

Black lines on coloured backround correspond to particles in chromosphere

Emission Spectra

Atoms/molecules/ions raised from ground state energy level when they absorb energy

Become excited

Lose energy by emitting EM radiation

Spectra of coloured lines on black backround

Continuous and Atomic Spectra

White light contains all visible wavelengths, and has a continuous spectra. Light from stars is not the same.

Atomic Spectra of Hydrogen Atoms

Balmer in visible light

Lyman in UV light

Spectra are the result of the interaction of light and matter

Bohr’s theory and Wave-Particle Duality (WPD)

Wave Theory of Light

Light is a form of EMR, so has wavelength and frequency

Moves at the speed of light in a vacuum

Different colours have different wavelengths

\[c=\lambda \times v\]

Particle Theory of Light

Light is a stream of packets of energy

Photons

\[E=h \times v\]

Bohr’s Theory

An excited atoms electrons will jump into higher energy levels

When they drop back, they emit EMR, giving an emission spectrum

When white light is passed through a cool sample of a gaseous element, black lines appear in the absorption spectrum. These correspond to the frequencies absorbed by the atoms in the sample. The intensities of these lines provide a measure of abundance.

EL3

Shells of Electrons

It is more appropriate to talk about shells than energy levels, due to complexity of atoms beyond hydrogen.
Each shell has a maximum number of electrons that it can hold. For a higher value of \(n\), the shell is further from the nucleus and has greater energy.

First Shell

\(n=1\)

2 electrons

Second/Third Shells

\(n=2\)/\(n=3\)

8 electrons

Fourth/Fifth Shells

\(n=4\)/\(n=5\)

18 electrons

Sixth/Seventh Shells

\(n=6\)/\(n=7\)

32 electrons

The lowest energy shells are filled first. Much of chemistry is decided by the outer shell electrons.

Sub-Shells of electrons

Sub-Shells are labelled s, p, d, f. These correspond to the shells:

\(n=1\) has an s sub-shell

\(n=2\) has s and p

\(n=3\) has s, p, and d

\(n=4\) has s, p, d, and f

Sub-Shell

Maximum number of electrons

s

2

p

6

d

10

f

14

In atoms other than hydrogen, sub-shells within a shell have different energies. The shells of 3d and 4s have an overlap in energies.

Atomic Orbitals

S sub-shells have one s-orbital

P sub-shells have three p-orbitals

D sub-shells have five d-orbitals

F sub-shells have seven f-orbitals

In an isolated atom, orbitals within the same sub-shell have the same energy.

Each orbital can hold a max of 2 electrons

Must have opposite spin

Corresponds to clockwise or anti-clockwise

The position of an electron is mapped with a probablilty function as it cannot be pinpointed exactly.

Filling Atomic Orbitals

The orbitals are filled to give the lowest energy arrangement possible. To do so, they are filled in order of increasing energy.